1. IIIIIIIV Chemical Bonding Introduction to Bonding

2. A. Vocabulary Chemical Bond attractive force between atoms or ions that binds them together as a unit

3. IONIC COVALENT Bond Formation Type of Structure Solubility in Water Electrical Conductivity Other Properties e - are transferred from metal to nonmetal high yes (solution or liquid) yes e - are shared between two nonmetals low no usually not Melting Point crystal lattice true molecules B. Types of Bonds Physical State solid liquid or gas odorous

4. “electron sea” METALLIC Bond Formation Type of Structure Solubility in Water Electrical Conductivity Other Properties Melting Point B. Types of Bonds Physical State e - are delocalized among metal atoms very high yes (any form) no malleable, ductile, lustrous solid

5. Ionic Bonding - Crystal Lattice RETURN B. Types of Bonds yRemember: Opposites Attract!

6. Covalent Bonding - True Molecules RETURN B. Types of Bonds Diatomic Molecule

7. Metallic Bonding - “Electron Sea” RETURN B. Types of Bonds

8. C. Bond Polarity  Difference in electronegativity determines bond type.  Above 1.7 = ionic  0.3-1.7 = polar covalent  0- up to 0.3 = non-polar covalent

9. C. Bond Polarity Polar Covalent Bond more e - neg atom has a partial negative charge   - less e - neg atom has a partial positive charge   +

10. Nonpolar Covalent Bond e - are shared equally usually between identical atoms Ex. F 2 C. Bond Polarity

11. ++ -- Polar Covalent Bond e - are shared unequally results in partial charges (dipole)

12. Nonpolar Polar Ionic View Bonding Animations.Bonding Animations C. Bond Polarity

13. Examples: zCl 2 zHCl zNaCl 3.0-3.0=0.0 Nonpolar 3.0-2.1=0.9 Polar 3.0-0.9=2.1 Ionic

14. D. Lewis Structures zElectron Dot Diagrams yshow valence e - as dots yEX: oxygen 2s2p O X

15. D. Lewis Structures zCovalent – show sharing of e - zIonic – show transfer of e -

16. D. Lewis Structures zCovalent – show sharing of e - zIonic – show transfer of e -

17. ++ -- ++ D. Lewis Structures zNonpolar Covalent - no charges zPolar Covalent - partial charges

18. 18 Steps for Building a Dot Structure for Covalent Compounds Ammonia, NH 3 1. Decide on the central atom; never H. Why? Most of the time, this is the least electronegative atom Therefore, N is central on this one 2. Add up the number of valence electrons that can be used. H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons NH 3

19. 19 3.Form a single bond between the central atom and each surrounding atom (each bond takes 2 electrons!) H H H N Building a Dot Structure H H H N 4. Remaining electrons form LONE PAIRS to complete the octet 3 BONDING PAIRS and 1 LONE PAIR.

20. 20 5.Check to make sure there are 8 electrons around each atom except H. H should only have 2 electrons.H should only have 2 electrons. Building a Dot Structure 6. Also, check the number of electrons in your drawing with the number of electrons from step 2. If you have more electrons in the drawing than in step 2, you must make double or triple bonds. If you have more electrons in the drawing than in step 2, you must make double or triple bonds. H H H N

21. 21 Carbon Dioxide, CO 2 1. Central atom = 2. Valence electrons = 3. Form bonds. 4. Place lone pairs on outer atoms. This leaves 12 electrons (6 pair). 5. Check to see that all atoms have 8 electrons around it except for H, which can have 2. C 4 e- O 6 e- X 2 O’s = 12 e- Total: 16 valence electrons

22. 22 Carbon Dioxide, CO 2 6. There are too many electrons in our drawing. We must form DOUBLE BONDS between C and O. Instead of sharing only 1 pair, a double bond shares 2 pairs. So one pair is taken away from each atom and replaced with another bond. C 4 e- O 6 e- X 2 O’s = 12 e- Total: 16 valence electrons How many are in the drawing?

23. 23 Double and even triple bonds are commonly observed for C, N, P, O, and S H 2 CO SO 3 C2F4C2F4C2F4C2F4

24. 24 Now You Try One! Draw Sulfur Dioxide, SO 2